Lab 5: Chemistry Unleashed



We’ve learned the terms, and now it’s time to explore dry ice, a plasma ball, and a few other cool chemistry devices and substances.

Lab Safety:

Even though we won’t have a formal lab write up this week, we will need to take laboratory safety very seriously.

For the dry ice:

  • Always use padded gloves (not the thin chemical kind) and tongs when moving dry ice. Never touch it with your bare skin. At -109° F, it can give you a severe burn.
  • Always use eye protection with dry ice.
  • You know not to eat or drink in lab, but particularly do not eat or drink anything with dry ice in it.
  • We will need to keep the dry ice outside the room so it doesn’t build up too much carbon dioxide gas in an enclosed space.

For the plasma ball and fluorescent light bulb:

  • Both these items are made out of glass. Please treat them gently and absolutely no horseplay!
  • We will be using eye protection.
  • The plasma ball can give you a burn to your fingertip if you place a metal item on it and then hold your finger over it.
  • Keep the plasma ball away from water and make sure your hands are dry.

For the acetone and Styrofoam:

  • Acetone is flammable. Keep it away from the plasma ball or any source of flame or sparks. Cell phones are potential sources of sparks.
  • We will also be using the acetone outside to reduce the chance of breathing fumes.
  • Please wear your eye protection.

For nitinol wire, be careful of the hot water and also keep your face away from the wire when you add it to the hot water. It can move violently and rapidly.

Part 1. Exploring solid carbon dioxide or dry ice.


  • Dry ice from the grocery store
  • Cooler (don’t close lid tightly)
  • Oven mitts and or heavy gloves
  • Metal tongs
  • Optional: metal knife or spoon
  • Quarter

Here are some suggestions for dry ice experiments (direct link):

If that isn’t enough, you can also make dry ice sing by placing a metal spoon on it (direct link).

Since you already have the dry ice, might as well try the next activity, too.

Part 2. Creating a dry ice and acetone cooling bath (-108° F).


  • Glass beaker
  • Dry ice
  • Acetone
  • Tongs
  • Heavy gloves

Put about 50 mL of acetone in a beaker and then slowly add golf ball or smaller-sized pieces of dry ice using gloves and/or tongs.

This “cool” Flickr video shows you the preparation and use of an acetone cooling bath.

Also, add Styrofoam to acetone.

Part 3. Experiments with a plasma ball.


  • Plasma ball
  • Extension cord
  • Fluorescent tube
  • Diffraction grating

Here are some good plasma ball demos (direct link):

Part 4. Nitonol Wire

  • Nitinol wire sample
  • 2 glass beakers (if have microwave)
  • Saucepan  to heat water (if have stove)
  • Ice
  • Water

This video might give you some ideas (direct link):

How about making a Nitinol Wire Inchworm (direct link)?

Lesson 5: Atoms and Elements

This week’s lesson is jam-packed with important information. Please take time to go over it carefully, and write down any questions you might have.

Textbook Reading:  Chapter 4, pp. 91-114




Atoms are pretty amazing to think about, and our knowledge of them is increasing all the time.

I have a friend who used to ask me, “Has anyone ever seen an atom?” Recently scientists have developed technology that allows us to do just that. It is called an atomic force microscope/scanning tunneling microscope.

This video from Nature shows us what atoms look like through one of these microscopes, as well as introducing some of the reasons why looking at atoms is useful (direct link).

How small are those atoms we just looked at? This TED video helps put it all into scale (direct link)

Finally, in this video Mr. Causey explains atoms, isotopes and ions and then helps you figure out how many protons, neutrons and electrons are found in some common examples (direct link).

Some other useful resources:

Theodore Gray has gone a long way towards making the elements more easy to relate to with his periodic table of element photographs at To look at each element, click on the photograph. His periodic table is also available as a book and as cards.

You might also want to check out this interactive period table.

Should you memorize the elements? I don’t think it is necessary to memorize every one, but you would benefit from learning the abbreviations for the most common elements, especially when we start doing reactions. Free Rice has a quiz/game to help you learn the basic symbols and once you’ve them, the full list of symbols. What level can you get to?

Optional video:
Are you interested in history? Bozeman Science gives an overview of how various parts of the atom were discovered and how the model of the atom changed. (direct link)

Since you are getting a break with lab this week, be sure to spend some extra time on this very important chapter. Feel free to contact me if you have any questions.

Sidebar: Recipes for Endothermic and Exothermic Reactions

This week we learned about endothermic and exothermic reactions in Lesson 4. Now let’s take a look at examples of these reactions.

Reaction 1.


  • 4 3.68 g packets Crystal Light Natural Pink Lemonade (primary ingredient citric acid)
  • Baking soda
  • Water
  • Large Styrofoam cup
  • Thermometer
  • Spoon or stirring rod

Procedure 1:

  1. In a sink, tub, or outdoors in an area that can get messy, mix the 4 packets of Crystal Light into approx. 100 mL of water in the Styrofoam cup.
  2. Take the temperature of the solution, taking care not to rest the thermometer on the bottom or side of the cup.
  3. Add about 1/3 cup baking soda.
  4. As the reaction starts to slow, take the temperature again.
  5. Adjust the amounts of ingredients if the temperature change was not apparent.

Did the temperature go up or down?



Reaction 2:


  • 2 Tablespoons active yeast
  • Water
  • Large Styrofoam cup
  • Thermometer
  • Hydrogen peroxide (from drug store)
  • Spoon or stirring rod

Procedure 2.

  1. In a sink, tub, or or outdoors in an area that can get messy, mix the yeast  into approx. 100 mL of water in the Styrofoam cup. (You may also add a squirt of dish detergent)
  2. Take the temperature of the solution, taking care not to rest the thermometer on the bottom or side of the cup.
  3. Add about 1/3 cup hydrogen peroxide.
  4. As the reaction starts to slow, take the temperature again.
  5. Adjust the amounts of ingredients if the temperature change was not apparent.

Did the temperature go down or up?

If you chose to, leave a comment to let us know how the experiment turned out.

Lab 4: Heat Capacity

Today we are going to investigate energy flow and specific heat capacity using a coffee cup calorimeter.

Experimental Title: Lab 4:  Heat Capacity

Date of laboratory:  June 24, 2014

Purpose: The purpose of this laboratory is to investigate the flow of thermal energy between substances until they reach thermal equilibrium.


Thermal energy is transferred from an object or substance with higher thermal energy to one at a lower thermal energy, until the two reach thermal equilibrium. 

Chemists often use a calorimeter to study thermal energy transfer. A common set up is called a “coffee cup calorimeter.” It consists of an insulated Styrofoam cup and a thermometer. To improve sensitivity, sometimes two Styrofoam cups are used, one inside the other and a cardboard, plastic or cork lid is added with a hole for the thermometer. This provides further insulation and less loss of thermal energy to the surroundings.

How much the temperature of a given substance increases for a given amount of heat transferred varies from substance to substance. The heat capacity is how much heat (measured in joules) is needed to raise the temperature of a given amount of substance 1°C. The specific heat capacity assumes the amount of the substance is measured in grams.

Equation:  q = m x C x ∆T

  • where q = amount of heat absorbed (joules)
  • m = mass (in g)
  • C = specific heat capacity (j/g °C)
  • ∆T = change in temperature (°C)

The specific heat capacity (C) for water is 4.184 J/g °C  and for copper is 0.385J/g °C.

Special safety concerns for Lab 4:

  • Today we will be wearing eye protection.
  • If boiling water spills on you, run cold water from the sink onto the area immediately. Don’t think, just run to the sink.
  • If the glass thermometer breaks, do not pick it up with your bare hands. Notify your instructor immediately.
  • Be sure to wash your hands when you are finished with this lab.


  • Eye protection
  • Stove
  • Oven mitts
  • Pan
  • Thermometer
  • Water
  • Ice
  • Styrofoam cups
  • Graduated cylinders
  • Table top scale
  • Strainer or colander
  • Pennies
  • Spoons
  • Clock
  • Calculator


Part 1:
1.    Pour 100 mL of water at room temperature into a large Styrofoam cup.  Insert the thermometer, taking care not to touch the walls of the cup. Determine the initial temperature of the water and note it in your notebook (see suggested data table below). Remove the thermometer from the cup.
2.    You will be adding 100 mL of boiling water to the cup. Using the initial temperature of the room temperature water, and assuming the temperature of boiling water to be 100°C, predict the final temperature of the mixture when the two samples are combined. Record your prediction in your notebook.
3.    Using a graduated cylinder, measure 100 mL of water. Pour the water into the pan provided and heat on the stove until it boils. Using oven mitts, carefully pour the 100 mL of boiling water into the Styrofoam cup with the room temperature water. If you want to stir the water, use a spoon and not a thermometer.
4. Take a temperature reading every minute until the mixed water reaches thermal equilibrium (when the temperature no longer changes, probably between four and six minutes). Record the final temperature in your notebook.

Part 2:
Repeat Part 1 using 75 mL room temperature water in step 1 and 225 mL of boiling water in step 3.

Part 3:
Repeat Part 1 using 225 mL of room temperature water in step 1 and using 75 mL of boiling water in step 3.

Part 4:
Repeat Part 1 using 100 mL of room temperature water for step 1 and 100 mL ice water for step 3.

Suggested Water Temperature Table (Parts 1-4)


Part 5:
1.    Weigh 35 pennies using the scale. Record the mass. (It should be close to 100 g.)
2.    Measure out an equal mass of room temperature water (remember 1 mL water is approx. 1 g). Pour the water into a large Styrofoam cup, insert the thermometer and record the temperature. Remove the thermometer.
3.    Cover the pennies with water in a pot and heat the water to boiling.
4.    While the pennies are heating, predict the final temperature that will result when the hot pennies (assume 100°C) are mixed with an equal mass of water in the cup. Record this value in your notebook.
5.    Drain the pennies in a colander over the sink to remove as much water as possible. Pour the hot pennies into the Styrofoam cup. Measure the final temperature as before and record in your notebook.

Part 6:
1.    Dry the 35 pennies and weigh again.
2.    Measure out an equal mass of room temperature water. Pour the water into a large Styrofoam cup, insert the thermometer and record the temperature. Remove the thermometer.
3.    Cover the pennies with ice water.
4.    While the pennies are cooling, predict the final temperature that will result when the cold pennies (assume near 0° C) are mixed with an equal mass of water in the cup. Record this value in your notebook.
5.    Drain the pennies to separate them from the ice and water and pour them into the Styrofoam cup. Measure the final temperature and record in your notebook.

Water Plus Pennies Temperature (Parts 4 and 5)



Now calculate the amount of heat (q) for water and pennies. We will work on this together in lab.


Once you have completed the six parts, sit down and write a sentence or two to explain the results of each part.


Record any thoughts you have about the experiments, including:

  • Why did we perform the experiments in Styrofoam cups?
  • Possible improvements to the procedures or how to tweak techniques
  • How the results differed from your expectations
  • Suggestions for other experiments
  • What key concepts you learned about heat capacity

We’ll go over the key concepts together at the end of lab.

Please leave a comment or send an e-mail if you have any questions before our meeting.

Lesson 4: Energy

What is energy? Why is it included in a chapter about matter? How are the two related?

Textbook Reading:  Chapter 3, part 2, pp. 66- 83



I. Energy

We use the word energy every day. We hear phrases like “energy crisis,” “running out of energy,” or “high energy bills.” Unfortunately, what we commonly mean by “energy” in these cases is some kind of fuel or other consumable resource.

In science, energy does not mean fuel. Instead, it is a more abstract concept:  “the capacity to do work.” It is not really anything concrete. It is not an object, but instead is a property that objects have.

Here’s what the famous physicist, Richard Feynman, had to say about energy:

There is a fact, or if you wish a law, governing all natural phenomena that are known to date. There is no exception to this law – it is exact so far as is known. The law is called the conservation of energy It says that there is a certain quantity, which we call energy, that does not change in the manifold changes which nature undergoes. That is a most abstract idea, because it is a mathematical principle; it says that there is a numerical quantity, which does not change when something happens. It is not a description of a mechanism, or anything concrete; it is just a strange fact that we can calculate some number and when we finish watching nature go through her tricks and calculate the number again, it is the same.

(Feynman, R. (1963).The Feynman Lectures on Physics. Book 1. New York: Addison-Wesley.)

What he is saying is that energy is a mathematical principle and no more. Isn’t that something to get your mind around?

A useful way to think about it is to think of various situations as “energy stores” rather than energy itself. These are situations with potential capacity to do work.

Some examples of energy stores:
1. Chemical (for example, alcohol + oxygen)

fireChemical energy stores can be used to move automobiles (internal combustion) and hot air balloons.

2. Kinetic (found in a moving object)


3. Gravitational (due to the position of an object in a gravitational field)

If you are sitting in a tree and drop a large rock, the gravitational energy will be transformed to kinetic energy. If your aim is good, you could drive in a stake with the dropped rock.

4. Elastic (for example, in a stretched rubber band or compressed spring)


5. Thermal (in a warm object)

coffee-heat-smallerThermal energy is important in determining the states of matter, as shown in this video (direct link).  Be sure to watch this one because the animations will help you visual the changes that occur between different states.


6. Magnetic (in magnets that are attracting or repelling )

johnny_automatic_magnet(Clipart from OpenClipArt)

Think of all the work a magnet can do.

7. Electrostatic (in two separated electric charges that are attracting, or repelling)


8.  Nuclear (released through radioactive decay, fission or fusion)

II. Introduction to Endothermic and Exothermic Reactions

As we study chemical reactions later in the course, we will find out that sometimes they release energy and sometimes they absorb energy. Our textbook is giving only a brief introduction in this chapter.

This video also gives an overview of endothermic and exothermic reactions (direct link).

III. Temperature versus Heat

Heat and temperature are further examples of vocabulary words with precise meanings in science that aren’t used as precisely in other contexts.

Temperature: A measure of a substance’s thermal energy (using a thermometer).

Heat: Thermal energy transfer or exchange between substances or objects.


Make sure you understand how to convert temperature units back and forth from Kelvin, Celsius and Fahrenheit, pp 71-73.

IV. Heat Capacity

We will be investigating heat capacity in lab this week, so pay particular attention to the terms heat capacity, specific heat capacity and the formula on page 75.

Maybe we’ll learn something that will help keep us cool.

Science Sidebar: Kinetic Sieving or Spontaneous Stratification

This week we are going to be separating mixtures using laboratory techniques. It turns out that there are geological processes that result in the separation of particles of different sizes in nature. Depending on the field of science you belong to, this is called kinetic sieving or spontaneous stratification.

This cool video shows how sand and colored sugar separate into bands while being poured (direct link). It also talks a bit about separation of mixtures at the end.

According to Gray, Shearer and Thornton, “Kinetic sieving is so efficient in dry granular flows that in small scale experiments a layer of 100% coarse grains develops on top of a layer of 100% fines with a sharp concentration jump between them.” That is amazing!

We have assembled the apparatus shown in the video so you can also give it a try in lab.

Reference: Time-dependent solutions for particle-size segregation in shallow granular avalanches.
1. J.M.N.T Gray,
2. M Shearer and
3. A.R Thornton
Proc. R. Soc. A 8 March 2006 vol. 462 no. 2067 947-972

Lab 3: Separation of Mixtures

In this lab we are going to investigate different ways to separate mixtures of substances by taking advantage of their physical properties. It will be inquiry-based, and you and your group will get to decide which techniques you want to use to separate a heterogeneous mixture of sand, poppy seeds, beans, salt and iron fillings.

Experimental Title: Lab 3 Separation of Mixtures

Date of laboratory:  June 17, 2014

Purpose: The purpose of this laboratory is to learn procedures and techniques used to separate mixtures of substances based on their physical properties.


A mixture is a combination of two or more substances in varying proportions. Scientists often need to separate mixtures into their components for analysis or to use in an experiment. It is possible to exploit differences in physical properties to separate substances from a mixture.

Special safety concerns for Lab 3:

  • If anything spills, please clean it up immediately with a paper towel and let your instructor know.
  • If glass breaks, do not pick it up with your bare hands. Notify your instructor immediately.
  • Be sure to wash your hands when you are finished with this lab


    • A heterogeneous mixture of beans, sand, salt, poppy seeds, and iron fillings
    • Forceps
    • Magnet
    • Plastic bag with tie
    • Beakers
    • Water
    • Graduated cylinders
    • Filter paper
    • Transfer pipette
    • Soda bottle filter
    • Soda bottle distillation apparatus (soda bottle cut in half, with lid)
    • Ice
    • Newspapers
    • Aluminum foil
    • Containers to hold separated substances
    • Plastic spoons


For this laboratory you and your group will decide which of the following procedures you will use to separate the mixture you receive. Keep in mind that you may need to repeat some procedures at different stages of the process. Go ahead and write the procedures below in your notebook now and you can refer to them by number as you use them. For example, “We added 25 mL of water, and then used procedure 4 to filter the mixture.” Be sure to write down the actual steps you use on lab day as well.

Procedure 1. Sorting or Manual Separation

Substances may differ in this size, shape and color. If the differences are large enough, it may be reasonable to simply pick out one of the substances with a pair of forceps and place it in a separate container.


For example, it would be relatively easy to pick the white jellybeans out of this mixture.

Procedure 2. Use a magnet to remove certain metals

Some substances (certain metals) are attracted to magnets and others are not. You can use this difference in magnetic properties to separate some mixtures.

If the metal fragments are small, place the magnet in a plastic bag and tie it shut. 


The baggie covering will make the metal fragments easier to remove from the magnet.


Drag the magnet over the mixture to attract magnetic metals. Brush the attracted metals off the plastic bag into a separate container.

Side note:  Have you ever tried this at the beach? Sand naturally contains iron fragments. In some places you might even be able to find small meteorites.

The following procedures require the addition of water to the mixture. Remember that if you add water, you might be dissolving some of the substances in the mixture. For example, salt or sugar dissolve into water making a solution.

Procedure 3. Decanting

It is possible to separate some substances based on differences in density. For example, oil floats on top of water. Add water, allow the substances to separate based on density, and then pour the upper layer with the less dense materials into a separate container.

If you have never done it, this video shows the standard way to decant in chemistry (direct link).

Procedure 4. Filtration

Filtration takes advantage of differences in particle size to separate mixtures. Generally filtration in chemistry involves special glassware, such as shown in figure 3.14 on page 65 in the textbook.

In this lab, we will use a large soda bottle cut in two, with the top inverted into the bottom.


Place the filter paper into the top of the soda bottle filter. Pour the liquid to be filtered through the filter. Larger particles will be trapped in the filter, and the liquid and smaller particles will pass through into the catchment container. Remove the filter and invert into a dish. Scrape off the solids with a spoon, if necessary.

Procedure 5. Evaporation

Both evaporation and its cousin, distillation, depend on differences in boiling points to separate materials. For example, with a solution of salt and water, the water has a lower boiling point. When heat is applied, the water boils away and the salt is left behind.

It is also possible to leave the solution in the sun for several days. The heat from the sun evaporates the water, again leaving the salt behind.

Evaporation involves applying energy to a solution in the form of heat, usually to remove water from a solution.

We will probably not be using evaporation today.

Procedure 6. Distillation

Distillation also takes advantage of differences in boiling point. In this case, the gas/vapor is captured again via condensation, rather than being allowed to escape into the air. A typical laboratory setup for distillation is shown in figure 3.13 on page 65 in the textbook.


We will set up a distillation apparatus from a soda bottle that has been cut in half.

Place the homogeneous solution in the bottom of the soda bottle. Place an empty glass in the center. Then invert the top of the soda bottle (with the cap left on) into the bottom half. Press down so it fits tightly and doesn’t allow gases to escape. Fill the top of the soda bottle with ice. Cover with newspaper (insulation) and then aluminum foil. Set in the sun.

The water should evaporate from the bottom, condense on the top and then run into the cup.


Many other separation procedures are possible in chemistry. For example, in a future lab we will be using chromatography to separate pigments in ink.


Once you have completed the separation, sit down and write a sentence or two to explain the results.


Record any thoughts you have about the experiments, including:

  • Possible improvements to the procedures or how to tweak techniques
  • How the results differed from your expectations
  • Suggestions for other experiments
  • What key concepts you learned about separating mixtures

We’ll go over the key concepts together at the end of lab.

If you would like to learn more, check this online chemistry lab from Phoenix College.

Please leave a comment or send an e-mail if you have any questions before our meeting.

Lesson 3: Matter

Now we move to the “nucleus” of chemistry:  matter!

Textbook Reading: Chapter 3, page 55 to top of page 66. We’ll save the Energy part of the chapter for next week.

Supplemental information:

As we learned in the first lesson, matter is anything that takes up space and has mass. In this chapter, Tro discusses some different ways to classify matter.

1. States of Matter

Even though matter can be found all over the Universe, you only find it in a few forms, called the “states” or sometimes “phases.” The textbook introduces you to three states of matter:  solid, liquid and gas (page 57). Did you know that there are other states of matter as well? Now people generally recognize at least five states of matter!


Plasma is widely considered to be the fourth state of matter. What is it? Plasma is a gas in which the atoms are ionized, meaning there are free negatively-charged electrons and positively-charged ions.

This video explains plasma (if the viewer doesn’t work, here’s a direct link).

Hopefully we’ll get to explore more about plasma in the future.

Bose-Einstein Condensates

If plasma is super high energy, then Bose-Einstein condensates are the exact opposites. Satyendra Bose and Albert Einstein predicted that matter would change state at temperatures approaching absolute zero way back in the 1920s, but no one was able to verify the existence of these condensates until 1995.

Here, one of the scientists who made a Bose-Einstein condensate explains what they are like (direct link):

Interested in learning more? The Chem4Kids website covers the basics of the five main states of matter in a particularly clear way.


2. Classifying Matter According to Composition

Matter can also be grouped according to what makes it up. If it is a pure substance, it contains only one kind of atom (an element), or molecule with different kinds of atoms (called a compound). In this case, the ratio of atoms in the compound is always the same.

Mixtures contain varying amounts of atoms in a combination of substances. If the mixture is uniform, it is called homogeneous. If you can point to something in the mix and say that substance is different from the rest, then it is called heterogeneous.

Pay particular attention to Figure 3.8 on page 59. We will be going over some concrete examples in class.

3. Chemical and Physical Properties

We will be investigating the physical properties of matter in our laboratory this week. Some physical properties are boiling point, density, and color. We’ll be learning a lot more about chemical properties as we progress through the book.

This video goes over the differences between chemical and physical changes in more detail if you’d like some clarification.

(Direct link if viewer is not working)

4. Law of Conservation of Mass

When French chemist Antione Lavoisier figured out that phylogiston was not part of combustion, he also realized that nothing was being created nor destroyed during chemical reactions. We might not always immediately realize where each substance is going in a reaction, but eventually we can track it down.

Pay particular attention to that tiny sidebar on page 65. We now know the law of conservation of mass is an oversimplification. It works with most chemical reactions, but in some nuclear reactions changes in mass do occur.

That’s it. Not so bad was it?

If you have any questions whatsoever, please leave a comment, send an e-mail or comment to the Yahoo group.

Lab 2 Density of Solids and Measurement Challenge

Great job with your laboratory notebooks last week. Keep up the good work!

Today we are going to practice taking measurements using the techniques we learned in the lesson, plus investigate the density of solids. You will learn to make accurate measurements, estimate to the proper level of certainty, and apply rules for significant figures in calculations.

Experimental Title: Lab 2 Density of Solids and Measurement Challenge

Date of laboratory:  June 10, 2014

Purpose: The purpose of this laboratory is to measure the volume, mass and density of solid substances.

Introduction:  How to measure using estimation.


How would you measure the red line in this example? To take a measurement with the ruler above it, first you would count the spaces between the large numbers. There are 10 spaces in the example, so each space is 1/10 of the distance between the black-labeled marks. If the black marks represent centimeters, then each smaller mark is 1 millimeter apart.

The green mark just before the red arrow is 9/10ths of the distance between 7 and 8, which is 7.9 cm (or 79 mm).  Previously you might have reported the answer as 7.9 cm. In chemistry, however, you want to get a more accurate reading of this measurement because the red line actually extends past 7.9 cm. How do you do this when there aren’t any markings? Try to visualize 10 steps in the space between the 7.9 and 8.0. The easiest one to visualize would be 5 steps (5/10) or halfway between.  It is pretty clear the red arrow is less than halfway, so the length of the red arrow is less than 7.95 cm.

Now estimate halfway between 7.9 and 7.95. That would be 7.925, but you can’t see that accurately. The arrow is very close to half of the first half.  So, you could record the length as 7.92 or 7.93 cm, either one would be correct.

In summary, the first two digits (7.9) are measured without any estimation. They make 2 significant figures (also called significant digits). The last digit is an educated estimate, but it does give us more accuracy. Therefore, it is counted as a third significant figure. By estimating, you are getting a little more accuracy than what the markings read.

Don’t worry, this will become easier with practice.

Important equations:

Density can be calculated using the formula:

density= mass (g)/volume(mL or cm3)

Volume of a cube is   V= S3 where S = length of an edge

Volume of a rectangular prism is  V =lwh  where l is the length of the base, w is its width and h is its height

The volume of a triangular prism is V = AH  where A = the area of the triangular base or 1/2bh and H = the height of the prism

Special safety concerns for Lab 2:

  • If anything spills, please clean it up immediately with a paper towel and let your instructor know.
  • If glass breaks, do not pick it up with your bare hands. Notify your instructor immediately.
  • Be sure to wash your hands when you are finished with this lab


  • Relational GeoSolid® blocks
  • Plastic blocks
  • Rubber stoppers
  • Sample of metal A
  • Sample of metal B
  • Sample of metal C
  • Sample rock
  • Water
  • Graduated cylinders
  • Table top scales
  • Transfer pipette
  • Rulers
  • Calculator


Note:  Today you can do the parts in any order, so go ahead to another part and come back to finish if you need to do so. No need to wait for materials.

Part 1. Determine the Volume of a CUBE and a Rectangular solid with a ruler

Last week we found the volume of a liquid using a graduated cylinder. This week we are going to find the volume of regularly-shaped objects by measuring and using mathematical formulas. Remember that a cubed centimeter is equal to 1 milliliter.

Procedure 1.

  1. Obtain a Relational GeoSolid® cube.
  2. Measure the length of a side in cm. Verify the shape is a cube by making sure the other sides are the same length.
  3. Record the length in your notebook.
  4. Using a calculator, calculate the volume using the formula V= S3.
  5. Record your answer using the correct number of significant figures.

cube-volume-table6. Repeat using the rectangular prism using the formula V =lwh.

prism-volume-table(Edit) 7. Now check the volumes you obtained by filling the Relational Geosolid® shapes with water.  Pour the water into a graduated cylinder and measure the volume.

Optional 1:  Obtain the solid triangular prism and calculate the volume using the formula V = AH where A = the area of the triangular base or 1/2bh and H = the height of the prism

triangular-prism-volume-tableWeigh the prism to obtain its mass. Measure the sides and calculate the volume. Now calculate the density. According to the text the density of glass is 2.6 g/cm3. Do you think the solid triangular prism is made of glass?

Let’s check to see if liquid volume is really equal to calculated volume.

8. Obtain a solid plastic cube (letter die)
9. Measure three sides to determine if it is a cube. If it is a true cube, then use the formula for volume V= S3 . If not, the use the formula for the volume V =lwh.
10. Record the side lengths and calculated volume.
11. Place 25 mL of water in a graduated cylinder. Carefully drop in the letter cube.
12. Record the final water level. Calculate the volume by subtracting 25 from the final level. How do the two volumes compare?


Part 2. Determine the Density of an Irregularly-shaped object Using Water Displacement

Do you remember the density video from Lab 1? In it the narrator explained how to figure out a volume of an irregularly-shaped solid object by immersing it in water in a graduated cylinder and recording the difference in water level.

Procedure 2.

1. Obtain a small rubber stopper and a graduated cylinder.
2. Weigh and record the mass of the dry stopper. Use the more accurate smaller scale.
3. Use tap water to fill your graduated cylinder to 25 mL.
4. Read and record this volume to the nearest 0.1 mL remembering to read the volume at the bottom of the meniscus.
5. Carefully submerge the rubber stopper in the graduated cylinder.
6. Read and record the new volume. What is the volume of the rubber stopper?
7. What is the density of the rubber stopper?

Leave room to record your observations in your notebook.

Part 3. Identify unknown samples using density

  • Sample of metal A
  • Sample of metal B
  • Sample of metal C
  • Sample rock

Use what you have learned in the part 2 of this lab to calculate the density of metal samples and find out their identity. Be sure to take careful notes of what you do and what your results are.

Repeat procedure 2, replacing the rubber stopper with metal samples.


Use this table of the density of some common substances to identify your unknowns.

Substance        Density (g/cm3)
Air                          0.0013
Wood (oak)          0.85
Water                     1.00
Ice                          0.93
Aluminum             2.7
Lead                        11.3
Copper                   8.96
Gold                        19.3
Iron                         7.86
Pyrite                       5.0
Galena                     7.5
Zinc                         7.133 – 7.14

(Edit) Optional 2. Measure the length, width and height of a wood block. Calculate the volume using V =lwh. Weigh the wood block. Now calculate its density. Is the wood block more or less dense than oak? Why is it preferable to calculate the volume of the wood block rather than use the water displacement method?

Final optional:  Obtain and weigh 5 dry pennies using the more accurate smaller scale. Now take their volume using the water method and calculate the density. Based on the density you obtained, do you think pennies are pure copper?

Please leave a comment or send an e-mail if you have any questions before our meeting.

Lesson 2: Measurement and Problem Solving

For this lesson we’re going to learn about scientific notation, significant figures, and units of measurement. If you have taken other science courses in the past, you are likely to find at least some of this section to be a review. For those of you who have never experienced these techniques and concepts, this is undoubtedly the most tedious section in the book. Give it your best effort, however, because once you’ve learned it, you will be able to apply it to many fields.

Textbook Reading:  Chapter 2 Measurement and Problem Solving (pp. 11-44)

Helpful practice:
Please do Skillbuilder 2.4 at the top of page 17 (answers on page 53).
Be sure to read and understand Examples 2.18 – 2.25 on pages 40-42. If you struggled with the density calculations in the last lab, look at examples 2.27 and 2.28 on pages 43-44.
Also, do Problem 42 on page 46 (answers for even numbered problems are in the back).


Still unsure about scientific notation after reading the text? Math is Fun has a scientific notation tutorial where you can type in your own examples to test (optional).

This video goes over significant figures (like in the textbook), and then gives a cool shortcut to use at the end.

(If the video player doesn’t work, link to YouTube)

Finally, this video gives a laid back review of measurement and units. You can zone out when he mentions accuracy, precision and percent error, as those are not covered in the Tro text.

(Direct link)

Please let me know if you have any questions. We will be going over examples at our meeting.

A Chemistry Sidebar:

mL vs ml

Have you seen milliliter abbreviated mL or ml and wondered which is correct?

According to the U.S. Metric Association (USMA):

“The symbol for liter (or litre) may be either a capital el (L) or a lowercase el (l); both are correct. In the U.S., Canada, and Australia, the capital el (L) is preferred, but most other nations use the lowercase el (l).”

So there you have it!

Atoms rule!